ELECTRODE  POTENTIAL  OF  MAGNESIUM 


BY 

RALPH  KENNY  HAMILTON 


THESIS 


lor  the 


DEGREE  OF  BACHELOR  OF  SCIENCE 

IN 

CHEMISTRY 


COLLEGE  OF  LIBERAL  ARTS  AND  SCIENCES 

UNIVERSITY  OF  ILLINOIS 


1921 


■ 


/ 9 2./ 

H I 82 


UNIVERSITY  OF  ILLINOIS 


Ju  rue.  _ JL i92_l__ 

THIS  IS  TO  CERTIFY  THAT  THE  THESIS  PREPARED  UNDER  MY  SUPERVISION  BY 

. j&lpJk  _ _Lenr.:ey_  _ _ xt  cn 

entitled Q tr  od  e_  _ _ P o ten_t  i &1  _ sl£  _ _M&_gOfis  iim. 


IS  APPROVED  BY  ME  AS  FULFILLING  THIS  PART  OF  THE  REQUIREMENTS  FOR  THE 
DEGREE  OF  


HEAD  OF  DEPARTMENT  OF 


Wo 
. ^ 


TABLE  OF  CONTENTS 


esarzosBEanwE 


Introduction 

Part  I Electrolytic  Potential  of 

Apparatus 

Historic  cil"*  . — - - — — * 

Experimental ---• 

Discussion  of  Results 

Part  II  Electrolytic  Potential  of 

Historical — 

Experimental--*--- — 

Discussion  of  Results ------ 

Acknowledgement * 


page  5 


Silver 


" 16 

” 19 

Magnesium 
-------  • " 2 1 

” 24 

--------  " 34 

36 


Digitized  by  the  Internet  Archive 
in  2016 


https://archive.org/details/electrodepotentiOOhami 


ILLUSTRATIONS 


Figure 

Figure 

Figure 

Figure 

Figure 

Figure 

Figure 


I Silver  Electrode  System 

II  Calomel  Electrode  Vessel 

III  Silver  Electrode  Vessel 

IV  Magnesium  Electrode  System 

V Hydrogen  Electrode 

VI  Magnesium  Electrode  and  Ves 

VII  Electrode  Potential  Curve- - 


„ 


INTRODUCTION 


The  experimental  determination  of  the  electrolytic  pot- 
ential of  the  elements  is  a field  in  which  a relatively  small 
amount  of  work  has  been  done  and  one  which  presents  a large 
opening  for  future  experimental  research.  On  account  of  the 
many  difficulties  in  the  way  of  making  significant  measure- 
ments, only  a very  few  metals  have  been  covered  with  any  de- 
gree of  certainty.  The  phenomena  of  passivity,  overvoltage, 
hydrolysis  of  the  sa,lts,  and  the  extreme  chemical  activity 
of  the  alkalie  and  alkaline  earth  metals  all  enter  in  until 
it  is  little  wonder  that  so  much  remains  undone.  Sven  in 
the  case  of  a metal  as  silver,  in  which  the  above  mentioned 
difficulties  do  not  present  themselves,  there  still  remains 
a considerable  lack  of  agreement  between  the  results  obtained 
by  different  experimentors . 

This  work  was  carried  out  with  the  purpose  in  view  of 
determining  the  potential  difference  between  metallic  mag- 
nesium and  a solution  of  magnesium  salt  normal  with  respect 
to  magnesium  ion.  This  potential  difference  had  not,  as  yet, 
been  determined  with  any  degree  of  accuracy.  Preliminary 
to  this  the  same  experiment  was  run  with  silver  using  some 
variations  in  experimental  procedure  which  had  been  successful! 


u s e d wi t h copper. 


PART  I 


ELECTROLYTIC  POTENTIAL  OF  SILVER 


A 


APPARATUS 


7 

APPARATUS 

Two  potentiometer  systems  uning  the  same  working  battery, 
standard  cell,  and  galvanometer  were  used.  The  working 
battery  was  a two  cell  lead  storage  battery  of  large  capacity. 
It  was  found  that  by  allowing  the  battery  to  remain  on  dis- 
charge, a current  of  one  mi Hi amp ere  being  continuously  with- 
drawn from  the  lead  cell,  a remarkably  constant  working  volt- 
age could  be  obtained,  decreasing  only  one  or  two  millivolts 
in  twenty-four  hours.  Further,  it  was  necessary  to  charge 
the  battery  only  once  in  one  to  two  months.  The  standard 
cell  was  a Weston  cell  of  voltage  1 .0136,  and  the  galvanometer 
a Leeds  and  Northrup  Type  K galvanometer.  The  figure  of 
merit  of  the  galvanometer  was  .343  X 10"'  amperes  per  milli- 
meter deflection. 

The  first  potentiometer  system  consisted  of  a,  Leeds  and 
Northrup  student’s  potentiom  ' : .3  ' y >f  ±.'J 

volts  and  the  second  a system  of  resistance  boxes  giving  an 
accuracy  of  ±. 1 millivolts.  This  combination  proved  to  be 
very  satisfactory.  By  means  of  the  student's  potentiometer 
a reading  to  ±1  millivolt  could  be  obtained  rapidly  while  with 
the  system  of  resistance  boxes,  readings  could  be  obtained 
to  the  next  decimal  place  which  was  in  keeping  with  the  con- 
stancy and  reproducibility  obtained  for  the  electrode  potential 
of  silver.  Dial  resistance  ooxes  were  found  very  convenient 
for  the  two  boxes  on  which  the  readings  were  made.  The  ac- 
companying blue  print,  Figure  I,  gives  a diagramatic  sketch 
of  the  set  up  as  used. 

The  potentiometer  system  is  the  same  as  is  ordinarily 


1 


' 


9 


used.  To  regulate  tine  working  current  to  one  mi  Hi  ampere, 
throw  switches  one  and  two,  set  the  potentiometer  at  1 .0136 
volts  and  use  tapping  key  (1).  Regula.te  resistance  R?  to 
zero  deflection  of  the  galvanometer.  For  the  unknown  electro- 
motive force,  throw  switches  one,  three  and  nine  and  use 
tapping  key  ( 1 ) . 

The  resistance  box  system  is  somewhat  more  complicated. 

The  working  current  in  this  case  is  one-tenth  milliampere. 

Box  R;  has  a resistance  of  from  1 to  9999  ohms  and  is  used 
to  regulate  the  working  current . Box  R^  has  11,000  ohms. 

To  regulate  the  working  current,  switch  five  is  thrown.  Box 
R^  has  a resistance  of  9313  ohms,  and  box  R7  has  10,186  ohms 
making  a total  of  19 ,999  ohms.  From  box  R7  the  circuit  re- 
turns to  the  negative  pole  of  the  lead  cell.  In  parallel 
with  box  R?  is  the  standard  cell  and  galvanometer.  Contact 
is  made  by  key  (2) . 

m determining  an  unknown  electromotive  force,  switch 
four  is  thrown.  When  the  unknown  electromotive  force  is 
between  zero  and  one  volt  it  is  placed  in  parallel  with  box 
Rs } that  is  switches  seven  and  eight  are  thrown  and  contact 
made  by  tapping  key  (3) . The  total  resistance  in  the  dial 
boxes  Rj  and  R^-  is  kept  at  9999  ohms  and  the  electromotive 
force  read  from  box  R^-,  one  ohm  corresponding  to  one-tenth 
millivolt.  When  the  unknown  electromotive  force  is  between 
one  and  two  volts,  switch  six  is  thrown  instead  of  seven. 

Figure  II  is  an  illustration  of  the  type  of  calomel  half- 
cell used.  In  the  reservoir  (A)  is  kept  a saturated  solution 
of  mercurous  chloride  in  tenth  normal  potassium  chloride  and 


1 1 

is  used  to  flush  out  the  arm  occasionally  to  do  away  with  the 
diffusion  of  the  stronger  potassium  chloride  solution  in  the 
salt  bridge  into  the  cell.  Figure  III  shows  the  type  of 
electrode  vessel  used  for  silver.  The  stop-cock  is  not  greased 
and  is  kept  closed  to  prevent  syphoning  over  of  the  solution. 


3 

HISTORICAL 


13 

HISTORICAL 

Although  a review  of  the  literature  on  the  electrode 
potential  of  silver  in  various  solutions  shows  that  a great 
deal  of  work  has  been  done  on  this  subject,  still  the  in- 
formation at  hand  from  which  the  electrolytic  potential  can 
be  calculated  is  very  small  indeed.  The  value  found  in  the 
older  literature  of  +.771  volts  as  determined  by  Willsraore, 
(Zeitschrift  ftir  physikalische  Chemie,  vol.  35,  pp.  2OI-332) 
must  certainly  be  in  error  by  as  much  as  .03  volt.  Of  the 
three  or  four  papers  bearing  directly  upon  the  subject,  that 
by  G.  N.  Lewis,  (Journal  of  the  American  Chemical  Society, 
vol.  23,  pp.  166-3)  is  probably  the  least  in  error.  His 
electrode  was  prepared  in  the  following  manner. 

A platinum  wire  was  sealed  into  a glass  tube  and  the 
projecting  part  wound  into  a spiral.  This  spiral  was  inserted 
into  a tube  containing  silver  oxide  and  the  whole  heated  at 
445eC.  until  all  the  oxide  was  decomposed.  The  platinum 
spiral  was  now  completely  enclosed  in  a loosely  cohering  mass 
of  finely  divided  silver.  The  cell, 

Pt:  Ag:  .IN  AgN03  : .IN  KN03  : .IN  KCl:  Hg^Cl,  : Hg 

was  found  to  give  more  constant  results  than  had  ever  previously 
been  obtained  for  silver.  This  type  of  electrode  is  undoubtedly 
superior  to  the  plated  electrode.  The  electrolytic  potential 
was  calculated  to  be  +.302  volts. 

Lewis  and  Lacey,  (Journal  of  the  American  Chemical  Soc- 
iety, vol.  36,  pp.  804-10)  working  on  the  copper  electrode 
have  found  that  copper  which  has  been  precipitated  electro- 
lytically  from  tenth  normal  copper  sulfate  solution  at  suf- 


14 


ficient  current  density  to  bring  down  a spongy,  non-adherent 
mass  of  metallic  copper  gave  a potential  which  was  far  more 
constant  than  could  be  obtained  by  copper  plating.  A copper 
plated  platinum  wire  was  used  to  make  contact  with  the  spongy 
copper . 

In  making  this  review  of  the  literature  on  electrode  pot- 
entials, No.  5 (1911)  and  No.  3 (1915)  of  Abhandlung  der 
deutschen  Bunsen  Gesellschaf t , on  Measurement  of  Electromotive 
Forces  of  Galvanic  Cells,  compiled,  by  Abegg,  Auerback  and 
Leuther  were  found  to  be  of  inestimable  service. 


c 

EXPERIMENTAL 


16 


EXPERIMENTAL 

Since  the  potential  of  even  the  most  promising  looking 

silver-plated  electrodes  show  considerable  variation,  it  was 
decided  to  prepare  a silver  electrode  analogous  to  Lewis' 
copper  electrode.  Tenth  normal  silver  nitrate  was  electro- 
lysed at  sufficient  current  density  to  bring  down  a dark, 
very  finely  divided,  spongy  precipitate  of  metallic  silver. 

The  silver  was  washed  with  tenth  normal  silver  nitrate  sol- 
ution and  allowed  to  stand  for  some  time  in  tenth  normal  silver 
nitrate  before  reading  the  potential.  Contact  was  made  by 
means  of  a very  bright  silver-plated  electrode.  The  cell, 

Ag:  .IN  AgNCL,  : KN03  sat.  sol:  .1  N.  E. 
gave  the  following  readings: 

Volts  Time  (min.) 

•3973  0 

.3977  5 

.3979  10 

.3973  65 

.3973  75 

This  cell  remained  constant  for  several  days.  However, 
readings  at  other  dilutions  did  not  give  the  desired  results. 
Further,  it  was  found  that  a silver-plated  platinum  wire  gave 
the  same  reading  in  a tenth  normal  silver  nitrate  solution 
whether  in  contact  with  the  finely  divided  silver  or  not. 

The  reading  is  this  case  was  higher  (.4002).  Also,  a smooth 
platinum  wire  placed  in  a tenth  normal  silver  nitrate  solution 
with  finely  divided  silver  gave  no  reading  at  all,  showing 
that  no  contact  was  made  between  the  platinum  and  the  finely 


• I) 


17 


divided  silver.  In  view  of  this  fact  it  was  decided  that 
black  amorphous  silver  obtained  at  high  current  density  was 
of  no  value  in  determining  the  electrode  potential  of  silver 
and  that  a smooth  platinum  wire  instead  of  one  plated  as  used 
by  Lewis  should  be  used  in  subsequent  experiments  since  the 
plated  wire  gives  a more  positive  potential. 

When  silver  is  deposited  at  a very  low  current  density  on 
a platinum  dish  small  white  crystals  of  silver  form.  These 
small  crystals  in  contact  with  a platinum  w ire  did  not  give 
satisfactory  results.  At  a slightly  higher  current  density, 
long  sharp  needles  of  crystalline  silver  form  in  clusters. 

It  was  this  silver  that  gave  all  that  could  be  desired  in  the 
way  of  constancy  and  reproduciabili ty . 

Working  with  freshly  prepared  solutions  of  silver  nitrate, 
the  following  cells  were  set  up  and  their  electromotive  force 
measured.  The  long  needle-like  crystals  of  bright  silver 
mentioned  above,  in  contact  with  a smooth  platinum  wire  was 
used  in  making  these  measurements.  Two  tenth  normal  calomel 
electrodes  were  prepared  and  their  potential  found  t,o  be  exactly 
the  same.  The  potentials  of  the  silver  electrodes  were  abso- 
lutely constant,  varying  at  no  time  by  more  than  one  tenth  milli- 
volt from  the  value  given.  All  readings  were  taken  at  22°  G . 


Ft: 

Cry st.  Ag:  AgN03 

.001N: 

KtJ03 

sat : 

. 1 N.  E.--- .2350 

If 

it  rt  ii 

.01N 

ft 

ft 

" -—.3415 

If 

n it  ti 

. IN 

ft 

II 

" --.3960 

■ 


■ 


. 


4 


D 

DISCUSSION  OF  RESULTS 


• ••  • 


19 


DISCUSSION  OF  RESULTS 

The  potential  of  the  tenth  normal  calomel  electrode  was 
taken  as  +-.336O  volts  on  the  hydrogen  scale.  With  the  aid 
of  conductivity  data  taken  from  Landolt  and  Bornstein,  the 
following  table  was  constructed. 


1 

2 

3 

4 

5 

6 

N 

Eq .Cond 

V.obs. 

E . P . Exp , 

. E.P.Calc. 

0.000 

1 15.3 

100.0 

0.001 

113.15 

97.7 

.2350 

.6210 

.6203 

0.01 

106.31 

93.1 

.3415 

.6775 

.6775 

0. 1 

94.33 

81.5 

.3960 

.7320 

.7322 

The  e 

lectrode 

potentials 

calculated  in  the  sixth  column 

made 

assuming 

the  value 

for  one- 

hundreth 

normal  to  be 

correct,  and  with  the  aid  of  the  Nernst  formula: 

E = .0530  log  c 
c' 

where  E is  the  difference  in  electrode  potential  and  c and 
c'  are  the  concentrations  of  the  silver  ion.  The  potential 
in  one-hundreth  normal  was  taken  as  probably  most  nearly  cor- 
rect because  at  greater  dilution  there  is  so  much  resistance 
in  the  cell  that  the  galvanometer  does  not  have  the  desired 
sensitivity  and  at  greater  concentration  the  thermodynamic 
environment  digresses  appreciably  from  that  of  pure  water, 
so  that  the  laws  of  ideal  solutions  would  not  be  obeyed. 

The  electrolytic  potential  calculated  from  the  value  obtained 
at  one-hundreth  normal  is  found  to  be  +.7953  volts.  This 
figure  shows  agreement  with  that  obtained  by  Lewis. 

A comparison  of  values  observed  in  column  five  and  those 


. 


20 


calculated  in  column  six  shows  an  agreement  more  close  than 
any  obtained  in  any  previously  published  work,  and  are,  in 
fact,  as  close  as  could  be  expected,  taking  into  consideration 
the  possible  error  in  alpha  as  obtained  from  the  conductivity 
data.  The  value  •+'.7953  volts  is,  therefore,  probably  the 
best  value  yet  obtained  for  the  electrolytic  potential  of  silver. 


part  ii 

ELECTROLYTIC  POTENTIAL  OF  MAGNESIUM 

A 

HISTORICAL 


21 


HISTORICAL 

The  electrode  potential  of  magnesium  has  never  been  det- 
ermined. with  any  decree  of  accuracy  on  account  of  the  form- 
ation of  basic  ss.lt s when  metallic  magnesium  is  placed  in 
contact  with  a magnesium  salt  solution,  (Kahlenberg,  Action 
of  Metallic  Magnesium,  Journal  of  the  American  Chemical  Soc- 
iety, vol.  25,  P*  3^0.)  Magnesium  attacks  water  slowly, 
hydrogen  being  evolved,  but  in  the  above  case  the  action  is 
much  more  rapid.  The  decomposition  is  probably  according 
to  the  following  equation: 

Mg  + MgCl,  + HOH—}  2 Mg ( OH)  Cl  -R  H^ 

It  is  worthy  of  note  at  this  point  that  magnesium  amalgams 
attack  ma,gnesium  salt  solutions  much  more  rapidly  than  does 
the  pure  magnesium.  M Le  Blanc,  (Zeitschrift  fiir  physikal- 
ische  Chernie,  vol.  5,  p.  467)  measured  the  voltage  of  the  cell: 
Mg)t(Hg):  MgCl^:  Znx(Hg) 

obtaining  a fairly  constant  electromotive  force  of  1.08  volts, 
but  it  was  our  experience,  as  will  be  shown  later,  that  such 
a cell  would  give  a rapidly  decreasing  electromotive  force. 

Neuman,  (Zeitscinrif  t fttr  physikali  sc  he  Chemie,  vol.  14, 
p.  215)  states  that  solid  magnesium  amalgam  decomposes  a mag- 
nesium salt  solution  very  rapidly  and  is  quickly  oxidized  by 
the  air.  The  potential  drop  between  magnesium  and  magnesium 
sulfate  was  found  to  decrease  with  the  time.  As  a mean  value 
he  gives  the  figure  -1.239  volts.  Comparisons  with  data  ob- 
tained in  this  laboratory  will  be  drawn  later. 

Kahlenberg,  in  his  article  previously  referred  to,  gives 
- I.366  volts  as  the  potential  of  magnesium  in  magnesium  sul- 


22 


fate.  This  value  was  found,  to  increase  with,  the  time.  Bab- 
orovsky,  (Zeitschrift  fttr  Electrochemie , vol.  11,  p.  465)  found 
-1.55  volts  to  be  the  potential  of  magnesium  in  magnesium  sul- 
fate and  that  the  potential  of  magnesium  amalgam  changed  very 
rapidly  with  the  time. 

The  work  of  Kistiakowsky,  (Zeitschrift  fiir  Electrochemie, 
vol.  14,  p.  113)  is  generally  accepted  as  being  the  most  sig- 
nificant. He  measured  the  potential  difference  between  a 
rotating  magnesium  electrode,  which  had  the  protruding  part 
covered  with  paraffin,  and  a normal  magnesium  sulfate  solution 
in  an  atmosphere  of  hydrogen.  He  took  the  view  that  it  was 
necessary  to  rotate  the  electrode  so  that  a homogeneous  sol- 
ution will  be  maintained  and  that  the  presence  of  oxygen  ef- 
fected the  potential  of  magnesium.  The  voltage  of  tne  cell: 
Mg:  MgSO^:  3N  KCl:  .1  N.  E. 

was  found  to  decrease  with  the  time,  the  highest  value  found 
being  1 .924  volts,  which  gradually  decreased  to  1 .790  volts 
at  the  end  of  an  hour  and  a half.  He  gives  the  highest  value 
found  as  the  most  nearly  correct.  A calculation  of  the  single 
potential  of  magnesium  by  Kelvin's  rule  gives  2.54  volts. 
Although  this  calculation  is  based  on  an  erroneous  assumption 
it  does  give  some  basis  for  comparison.  Thus,  the  individual 
experimentors  disagree  between  themselves  by  as  much  as  three 
tenths  of  a volt  and  on  comparing  experimental  values  with 
the  above  calculated  value,  one  would  say  that  probably  all 
of  them  are  far  from  the  correct  one.  In  view  of  this  state 
of  disagreement  it  was  decided  to  investigate  the  subject. 


B 

EXPERIMENTAL 


24 


EXPERIMENTAL 

When  experimental  difficulties  in  determining  the  elec- 
trode potential  of  an  element  are  involved,  as  is  the  case  with 
magnesium,  indirect  methods  have  been  employed  making  use  of 
organic  solvents,  (Marquis,  Electrode  Potential  of  Arsenic, 
Journal  of  the  American  Chemical  Society,  vol.  42,  p.  1 5 69 
and  Lewis,  Electrode  Potential  of  Sodium,  Journal  of  the 
American  Chemical  Society,  vol.  32,  p.  1459.)  It  is  quite 
evident  that  such  a solvent,  in  addition  to  dissolving  a mag- 
nesium salt,  must  cause  it  to  ionize  appreciably  into  the  pos- 
itively charged  magnesium  cation,  and  must  not  form  addition 
compounds;  nor  can  any  phenomena  similar  to  hydrolysis  be  pre- 
sent. On  recalling  the  G-rignard  reaction  one  'would  say  offhand 
that  a suitable  solvent  would  be  hard  to  find  and  such  was 
found  to  be  the  case.  Attempts  were  made  to  deposite  magnesium 

electrolyticaliy  from  solutions  of  magnesium  chloride,  magnesium 
sulfate,  or  magnesium  bromide  in  acetone,  methyl  alcohol,  or 
pyridine.  In  any  case,  magnesium  was  found  to  go  into  sol- 
ution at  the  anode  but  only  a gas  could  be  obtained  at  the 
cathode,  as  is  the  case  in  aqueous  solutions.  The  possibility 
of  using  an  organic  solvent  was  then  discarded. 

Kahlenberg,  in  his  investigations  of  the  action  of  met- 
allic magnesium,  states  that  alkalie  solution  are  not  attacked 
by  metallic  magnesium,  the  metall  remaining  bright  for  several 
hours.  At  Dr.  Reedy's  suggestion  the  following  cell  was  set  up: 
Mg:  Mg(OH)  ' .IN  KOH:  3N  KC1:  .1  N . E. 

Theoretically  this  seemed  quite  promising  for,  even  if  decomp- 
osition did  take  place,  all  that  could  occur  would  be  the  for- 


. 


. 


25 


ation  of  more  of  the  hydroxide.  It  was  found  that  although 
metallic  magnesium  does  not  decompose  potassium  hydroxide  sol- 
ution, hydrogen  is  evolved  when  magnesium  hydroxide  is  present 
in  the  solution.  Further,  a constant  electromotive  force 
could  not  be  obtained.  The  average  was  found  to  be  -I.O3 
volts  for  the  potential  of  the  magnesium  half-cell,  which  is 
about  the  same  as  has  already  been  obtained  for  magnesium  in 
potassium  hydroxide  solution  and  to  which  no  significance  can 
be  attached. 

Lewis,  (Journal  of  the  American  Chemical  Society,  vol . 

32,  p.  1459)  has  obtained  excellent  results  for  the  alkalie 
metals  by  making  use  of  dilute  amalgams.  As  was  previously 
mentioned,  M Le  Blanc  and  Neuman  did  not  agree  on  the  matter 
of  the  electromotive  force  of  magnesium  amalgams,  so  it  was 
decided  to  make  an  investigation.  An  amalgam  was  made  up 
containing  about  .5$  of  magnesium,  but  it  was  rapidly  oxidized 
by  the  air  forming  a solid.  A more  dilute  amalgam  was  made 
up  then  ( .01$)  and  it  too  was  rapidly  oxidized.  When  placed 
in  a magnesium  chloride  solution  a large  black  precipitate 
was  found  to  be  instantly  formed.  An  amalgam  containing  .001$ 
magnesium  gave  this  same  black  precipitate  when  in  contact 
with  magnesium  chloride  but  not  so  rapidly.  Its  potential 
was  measured  and  found  to  change  rapidly  with  the  time  as  is 
shown  by  the  following  taible . 


26 


Time 

Potential 

1 min . 

-1 .298  volt; 

4 " 

-1.332 

7 " 

-1 .329 

16  " 

-1.297 

33  ” 

- 1 .220  " 

56  " 

-0.364 

With,  this  the  possibility  of  making  use  of  a magnesium  amalgam 
was  given  up. 

A rotating  magnesium  electrode  similar  to  that  used  by 
Kistiakowsky  was  next  assembled  with  the  idea  in  mind  of  as- 
certaining whether  Kistiakowsky ’ s results  were  reproducible 
or  not.  The  metal  used  was  in  a circular  stick  about  one- 
half  inch  in  diameter  and.  three  inches  long.  Several  meas- 
urements were  made  in  an  atmosphere  of  hydrogen  but  the  result 
v/ere  the  same  in  either  case.  The  following  tables  give  a 
comparison  between  the  best  set  of  values  obtained  and  those 
given  by  Kistiakowsky.  The  cell  measured  in  either  case  is, 
Mg:  MgSO  : 3N  KC1:  .1  N.  E. 

the  only  difference  being  that  we  used  tenth  normal  magnesium 
sulfate  where  Kistiakowsky  used  normal.  Our  voltage  should 
therefore  be  about  .025  volts  higher  than  his. 

An  attempt  to  extrapolate  our  readings  back  to  zero  time 
was  a failure  because  it  was  impossible  to  connect  the  points 
by  any  kind  of  a curve.  However,  the  two  sets  of  readings 
do  agree  that  the  potential  of  the  magnesium  electrode  de- 
creases with  the  time.  This  is  caused  by  the  formation  of 


II 


27 


Kistiakowsky  This  laboratory 


Voltage 

Time 

Voltage 

Time 

1 .924 

0 

1 .379 

0:34 

1 .922 

1 

1 . 354 

1:13 

1 .916 

5 

1 .312 

1 : 56 

1 .900 

9 

1 .793 

2:18 

1 .900 

13 

1 .735 

2:34 

1 .376 

32 

1 .777 

2:56 

1 .346 

77 

1 .753 

3:38 

1 .790 

37 

1.631 

123 : 00 

the  basic  salt  which 

appears  on 

the  electrode 

immediately  on 

immersing  it  in  the 

solution. 

This  salt  will  dissolve  in 

hydrochloric  acid. 

The  electrode  potential 

seemed  to  reach 

a minimum  at  about  the  lowest  value  given,  this  minimum 
varying  somewhat  with  the  salt  used.  It  is  about  the  same 
as  that  obtained  by  Kahlenberg  but  could  hardly  be  accepted 
as  the  true  potential  difference  between  magnesium  and  mag- 
nesium salt  solution. 

Since  this  basic  salt  is  soluble  in  an  acid,  the  possi- 
bility of  measuring  the  electrode  potential  in  acid  solution 
suggested  itself.  Although  the  acid  would  dissolve  the  mag- 
nesium, it  was  thought  that  by  mixing  half  and  half  solution 
of  magnesium  sulfate  and  sulfuric  acid  of  the  same  strength 
and  obtaining  the  initial  reading,  a value  to  which  some  sig- 
nificance could  be  attaches  might  be  obtained.  However,  the 
electrode  potential  was  found  to  change  very  rapidly  with  the 
time  and  it  was  impossible  to  take  a reading  quickly  enough. 


...  ,« 


28 


The  final  method  used,  which  after  taking  a good  many 
readings  with  different  alterations  gave  some  degree  of  success, 
was  to  measure  the  electrode  potential  in  acid  solution  and 
simultaneously  follow  the  hydrogen  ion  concentration  by  means 
of  a hydrogen  electrode. 

The  first  theory  was  to  start  out  with  a tenth  normal 
solution  of  sulfuric  acid  and  let  it  react  with  the  magnesium 
until  completion,  when  the  solution  will  be  tenth  normal  mag- 
nesium sulfate.  The  end  point  was  to  be  determined  by  the 
hydrogen  electrode  which  would  show  a potential  of  .406  volts. 
Theoretically,  the  basic  salt  should  not  commence  to  form  until 
all  the  acid  was  gone  but  actually  it  did  not  wait  that  long 
so  that  this  method  had  to  be  modified. 

When  the  solution  is  tenth  normal  acid,  the  action  of 
the  acid  is  strong  enough  to  prevent  the  formation  of  the  basic 
salt,  so  it  was  decided  to  start  out  with  two  tenths  normal 
sulfuric  acid  and  let  the  reaction  run  past  the  half  way  mark, 
again  making  use  of  the  hydrogen  electrode.  Figure  IV  shows 
diagrammatic ally  the  simple  arrangement  make  use  of.  To  bal- 
ance the  system,  throw  switch  one;  to  measure  the  hydrogen 
electrode,  switches  two  and  three;  the  magnesium  electrode, 
switches  two  and  four.  Figure  V shows  the  kind  of  hydrogen 
electrode  used,  and  Figure  VI  the  electrode  vessel. 

The  voltage  of  the  hydrogen-calomel  cell  at  that  point 
where  the  two  tenths  normal  sulfuric  acid  has  been  changed  to 
a mixture  of  tenth  normal  magnesium  sulfate  and  tenth  normal 
sulfuric  acid  was  previously  found  to  be  .4075.  On  account  of 


31 


it  being  impossible  to  take  both  readings  at  tbe  same  time,  it 
was  necessary  to  take  readings  alternately  and  plot  them  against 
time  on  tbe  same  paper.  Obviously,  tbe  potential  of  tbe  mag- 
nesium electrode  at  that  time  when  tbe  potential  of  tbe  hydro- 
gen electrode  is  .4075  against  tbe  calomel  is  tbe  one  sought. 

A good  many  readings  were  taken,  four  sets  of  wnicb  are  given 
in  tbe  following  table. 


Table  I 

Table  II 

Mg 

Time 

H 

Mg 

Time 

H 

2.036 

0:00 

2.090 

0:00 

0:42 

.399 

0:40 

.400 

2.081 

1:37 

2.036 

1 : 10 

2:20 

.4077 

1 : 40 

.405 

2.069 

3:15 

2.030 

2:  15 

4:  10 

.4137 

2:53 

.4125 

2.053 

4:40 

2.067 

3:40 

Table  III 

laole  xV 

Mg 

Time 

H 

Mg 

Time 

H 

2.039 

0:00 

2.0345 

0:00 

0 : 43 

.397 

0:55 

.397 

2 . 035 

1 :25 

2 . 034 

1:18 

1:53 

.4032 

1 :53 

.401 

2.079 

2:34 

2.079 

2:33 

3:09 

.41  1 

04 

O 

O 

.4065 

2.0665 

3:48 

2.073 

303 

4:00 

.41  15 

c 

DISCUSSION  OP  RESULTS 


4 


34 


DISCUSSION  OF  RESULTS 

Two  graphs  for  each  set  of  data  were  drawn  on  a single 
sheet,  one  of  the  hydrogen  electrode  against  time  and  the  other 
of  the  magnesium  electrode  against  time.  From  this  the  pot- 
ential of  the  magnesium  electrode  at  the  same  time  that  the 
potential  of  the  hydrogen  electrode  showed  that  one  half  of  the 
acid  had  been  used  up  could  be  determined.  Figure  VII  was 
drawn  from  the  data  given  on  Table  III,  which  is  about  an  average 
of  all  four.  The  values  found  in  each  case  were: 

Table  I 2.0765 

" II  2.0307 

" III  2.0733 

" iv  2.0730 

Average  2 . 077 

This  gives  -1.741  volts  as  the  potential  of  magnesium  against 
the  solution  used. 

The  solution  is  a tenth  normal  solution  of  magnesium  sul- 
fate and  tenth  normal  sulfuric  acid.  Tenth  normal  magnesium 
sulfate  i3  44.1^  dissociated  and  tenth  normal  sulfuric  acid 
dissociated  so  that  it  contains  .0625  moles  hydrogen  ion,  . O315 
moles  bisulfate  ion,  and  *0155  moles  sulfate  ion  per  liter. 
Applying  the  isohydric  principle,  the  concentration  of  magnesium 
ion  is  found  to  be  .0258  normal.  By  means  of  the  Nernst 
equation,  the  electrolytic  potential  is  calculated  to  be  -1.700 
volts.  On  account  of  the  difficulty  in  determining  exactly 
the  concentration  of  magnesium  ion,  this  figure  may  be  in  error 
by  as  much  as  ±*015  volts. 


. 


. 


* 


ACKNOWLEDGEMENT 


36 


ACKNOWLEDGEMENT 

The  author  wishes  to  take  advantage  of  this  opportunity 
to  express  his  appreciation  to  Dr.  Dietrichson,  under  whose 
direction  this  work  has  been  carried  out.  To  his  many  help- 
ful suggestions  and  frequent  visits  to  the  laboratory,  which 
have  caused  this  work  to  be  a sourse  of  much  personal  pleasure 
and  satisfaction,  may  be  attributed  the  success  attained  in 
this  investigation. 


